4.01

Preface and Context to Hydrogen and Fuel Cells

AJ Cruden, University of Strathclyde, Glasgow, UK
© 2012 Elsevier Ltd.

4.01.1
4.01.2
4.01.2.1
4.01.2.2
4.01.2.3
4.01.2.4
4.01.2.5
4.01.2.6
4.01.2.7
4.01.2.8
4.01.2.9
4.01.2.10
4.01.2.11
4.01.2.12
4.01.2.13
4.01.2.14
4.01.3
4.01.3.1
4.01.3.2
4.01.3.3
4.01.4
4.01.4.1
4.01.4.2
4.01.4.3

4.01.4.4
4.01.5
References

Introduction
An Overview of This Volume
Chapter 4.01: Introduction
Chapter 4.02: Current Perspective on Hydrogen and Fuel Cells
Chapter 4.03: Hydrogen Economics and Policy
Chapter 4.04: Hydrogen Safety Engineering: The State-of-the-Art and Future Progress
Chapter 4.05: Hydrogen Storage: Compressed Gas
Chapter 4.06: Hydrogen Storage: Liquid and Chemical
Chapter 4.07: Alkaline Fuel Cells: Theory and Application
Chapter 4.08: PEM Fuel Cells: Applications
Chapter 4.09: Molten Carbonate Fuel Cells: Theory and Application
Chapter 4.10: Solid Oxide Fuel Cells: Theory and Materials
Chapter 4.11: Biological and Microbial Fuel Cells
Chapter 4.12: Hydrogen and Fuel Cells in Transport
Chapter 4.13: H2 and Fuel Cells as Controlled Renewables: FC Power Electronics
Chapter 4.14: Future Perspective on Hydrogen and Fuel Cells
Hydrogen and Fuel Cell Technology – Supplementary Material
Flow Cells or Regenerative Fuel Cells
Hydrogen Production – Electrolysis
Hydrogen Demonstration Units – State of the Art
Introduction to Basic Electrochemistry
Redox Reactions
Electrochemical Series
Gibbs Energy – Useful Work
Practical Fuel Cells
Conclusions


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4.01.1 Introduction
Hydrogen is the most abundant material in the Universe, forming over 75% of known matter; however, it does not commonly exist
on Earth in its natural form, due to its highly reactive nature, but within other compounds, most notably water and hydrocarbons.
The discovery of hydrogen gas is credited to the famous English philosopher Henry Cavendish (although he was actually born in

Nice, France!) who, in 1766, wrote a seminal paper entitled ‘Experiments on Factitious Airs’ [1] (Figure 1) after experiments
dissolving different metals (such as zinc) in acidic solutions.
These experiments produced a gas, the ‘factitious air’ that “takes fire, and goes off with an explosion” (see a further exert from the
Cavendish paper of 1766 shown in Figure 2), which is now a common high-school test for hydrogen gas – set fire to it and it goes ‘pop’!
Cavendish went on to determine that this gas was significantly lighter than air and, although credited with isolating this new
inflammable gas, it was another Frenchman, Antoine Lavoisier, who named this gas as hydrogen in 1783. Indeed, the name
‘hydrogen’ itself is from the Greek words ‘hydros’ (meaning ‘water’) and ‘generos’ (meaning ‘to make’ or ‘to create’); hence, the name
hydrogen means ‘to make water’ or ‘water former’.
The story of hydrogen took a further step forward around this time when the Englishman, William Nicholson, correctly
identified it following his early experiments on electrolysis. Figure 3 shows an extract of Nicholson’s famous paper of 1800 [2]
where he correctly determines that water is composed of hydrogen and oxygen.
Of course, the discovery and naming of hydrogen at this time is all the more challenging due to its properties which, at standard
temperature and pressure (STP), render it odorless, colorless, tasteless, nontoxic yet highly flammable (within its flammability
limits of 4–74% in air). It is a highly reactive substance (hence, it does not naturally occur but is found bonded within many other
compounds) and is the lightest element in the periodic table.
Hydrogen at STP is in the form of a molecular gas. It was not until 1898 that the Scotsman, Sir James Dewar, liquefied hydrogen
for the first time (see Figure 4 showing a repeat of this first experiment in 1899), achieving temperatures of 20 K or –253 °C. Even at
such extreme low temperatures, hydrogen formed a colorless liquid.
Dewar continued his pursuit of ever colder temperatures and was the first to produce solid hydrogen, at temperatures below 14 K
(–259 °C) in 1899.

Comprehensive Renewable Energy, Volume 4

doi:10.1016/B978-0-08-087872-0.00401-7

1


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Preface and Context to Hydrogen and Fuel Cells

Figure 1 Cavendish’s paper on ‘factitious air’ from 1766. From />
Figure 2 Extract from Cavendish’s 1766 philosophical transactions paper. />PA141&dq=philosophical+transactions+of+the+royal+society+1766+cavendish&source=bl&ots=IGBnzpS7_e&sig=5cTcFhZXXJQZLHV7Nf6NVx3sJfo
&hl=en&ei=alO6TsywKMi3hAfinZjBBw&sa=X&oi=book_result&ct=result&resnum=7&ved=0CEQQ6AEwBg#v=onepage&q=philosophical%
20transactions%20of%20the%20royal%20society%201766%20cavendish&f=false

From around this time, hydrogen gas had been in use as a constituent of coal gas. Coal gas, as opposed to natural gas (naturally
occurring gas containing methane as the majority component), was a manufactured gas from the ‘cracking’ of coal, that is, the
breakup of the long coal hydrocarbon molecules into different compounds. This cracking was achieved by controlled combustion
(with limited air) to produce a gas containing up to 50% hydrogen, methane, carbon monoxide, and other gaseous elements. This
gas was produced locally and used for lighting and heat/cooking [3].
The manufacture of coal gas in the United Kingdom decreased rapidly following the discovery and extraction of natural gas,
principally methane, from the North Sea, which could be both extracted and used in a much environmentally clean fashion, with
less remediation required, than coal gas production [4].
The use of coal gas was widespread and very visible in many towns and cities, principally due to the large gas holders (or
gasometers as they became known as) that stored the coal gas. These tanks (an example is shown in Figure 5) were a
common site and are only now being replaced by high pressure storage in modern, underground, plastic high pressure
natural gas pipelines.
So hydrogen gas has been in use, albeit in a mixed dilute form within a coal gas mix, for well over a century. Hydrogen
has also many industrial and speciality uses: as a product in semiconductor processing, petroleum refining, ammonia
production for fertilizer production, heat treatment of metals, as a coolant in large electrical generators in power stations,
and as a rocket fuel for space missions! However, this volume will concentrate on the technologies that aim to use pure
hydrogen as a fuel.
Hydrogen is not a source of ‘primary energy’, as hydrogen requires to be produced/released or manufactured as a pure gas, and
also requires further treatment to liberate energy when being converted to useful work. It is also not a form of renewable energy,


Preface and Context to Hydrogen and Fuel Cells


3

Figure 3 Extract from Nicholson’s paper of 1800, determining the composition of water. />id=TggAAAAAMAAJ&printsec=frontcover&source=gbs_ge_summary_r&cad=0#v=onepage&q&f=false

Figure 4 Painting of Sir James Dewar demonstrating liquefaction of hydrogen to the Royal Institution, 1899.

rather it should be viewed as an ‘energy vector’, in a similar fashion to electricity which does not naturally occur, and requires to be
produced (generated) from primary energy sources/fuels, and is then reconverted to useful work in our electric lights, heaters,
machinery, computers, and so on. However, hydrogen is a unique type of energy vector in that it can be stored, in large volumes,
unlike electricity, and this unique feature will enable its use to support development and implementation of the other forms of
renewable energy reported in this Comprehensive Renewable Energy series.


4

Preface and Context to Hydrogen and Fuel Cells

Figure 5 Coal gas storage tank (gasometer). />
This capacity for hydrogen to be used as both an energy store and a fuel is particularly relevant to the transport sector, a sector
dominated by fossil fuels and hence carbon emission concerns, and a sector that forms of renewable energy like wind or solar
(unless biofuels) is not frequently linked to. The use of hydrogen as an intermediate energy store and fuel will allow wind energy
and other forms of renewables to power water electrolysis plants, to produce hydrogen gas for use in vehicles, thereby creating a
‘double benefit’: increase the usage of renewable energy and replace vehicle fossil fuel consumption with a zero emission
alternative.
A significant contributory factor in the pursuit of hydrogen use as a fuel has been the parallel development of fuel cells, a DC
electrical source that can be thought of as a ‘continuously operating battery’. The fuel cell was first discovered by William Grove in
1839 [5], and was reported as a postscript in a paper focussing on developments on ‘voltaic series’, at the time an area of feverish
scientific development in the new field of electricity, following Alessandro Volta’s creation of the zinc/copper pile battery. Grove’s
postscript to his paper is shown in Figure 6, and marks the time when Grove, in a subsequent publication in 1842 [6], presented
further details of this initial fuel cell and included an image of the experimental arrangement, an image which has become redolent

with the history of fuel cells. A full copy of this paper is shown in Figure 7, first to illustrate the image of Grove’s fuel cell and second
to highlight the spirit of research at this time. (Note the comments on p. 418 regarding the detection of electric potential by means
of an electric shock felt across five persons joining hands!)
A principal difference between a ‘battery’ and a ‘fuel cell’ is that a battery contains (or holds internally) the ‘fuel’ or chemical
compounds with which to generate electricity within the casing of the battery: by contrast, the ‘fuel’ for a fuel cell is both held and
supplied externally, and hence can be continuously supplied or replenished if desired.

4.01.2 An Overview of This Volume
This volume, entitled ‘Hydrogen and Fuel Cell Technology’, part of the overall Comprehensive Renewable Energy Major Reference
Work, covers a range of technologies spanning the hydrogen and fuel cell sector. This particular chapter presents the context to this


Preface and Context to Hydrogen and Fuel Cells

5

Figure 6 William Grove’s postscript describing the first recorded operation of a hydrogen fuel cell.

sector and an overview of hydrogen and fuel cell technology. The remaining chapters of this volume will now be briefly introduced,
in the order they are presented within the volume:

4.01.2.1

Chapter 4.01: Introduction

The history and background to the production and use of hydrogen as a fuel within fuel cells is presented in this chapter, in addition
to providing an introduction to the fundamental thermodynamics of how fuel cells operate. This is as a precursor to the further,
more technically detailed, chapters that follow and includes an overview and context to the hydrogen and fuel cell sector. It also



6

Preface and Context to Hydrogen and Fuel Cells

Figure 7 Grove’s ‘Gaseous Voltaic Battery’ – the first reported hydrogen fuel cell.

presents information on examples of existing hydrogen production technology, namely, large-scale alkaline electrolysis, and of one
current transportable proton exchange membrane (PEM)-based hydrogen production system for vehicle refuelling. Finally, some
detail of a regenerative fuel cell, a novel type of fuel cell that is gaining traction as a potential energy storage technology, is presented
to illustrate that some fuel cell technologies do not require hydrogen gas to operate.


Preface and Context to Hydrogen and Fuel Cells

4.01.2.2

7

Chapter 4.02: Current Perspective on Hydrogen and Fuel Cells

This chapter presents a historical perspective and the current state of the utilization of hydrogen as a fuel, and of fuel cells as a power
source. Given the impetus to develop and use hydrogen and fuel cells within the space exploration sector, this chapter gives a
particular perspective of these early developments and an overview of current uses.

4.01.2.3

Chapter 4.03: Hydrogen Economics and Policy

Developing hydrogen gas as a new form of energy vector requires significant technical development, as later chapters within this
volume will discuss; however, it will also require significant financial investment. Additionally, to break into the classic ‘chicken and

egg’ cycle that tends to inhibit new technology uptake (e.g., users will not purchase hydrogen vehicles until there is an established
network of hydrogen refuelling stations, while energy companies will not develop or invest in a network of hydrogen refuelling
stations until there is a significant population of hydrogen vehicles to utilize this investment), the nature and use of government
policy instruments to stimulate this development is presented and discussed.

4.01.2.4

Chapter 4.04: Hydrogen Safety Engineering: The State-of-the-Art and Future Progress

A key issue surrounding the utilization of hydrogen as a fuel centers on the safe use of this highly reactive, flammable gas. The
subject of hydrogen safety is simultaneously both technical and emotive, with considerable ongoing technical work studying
hydrogen flammability and safety equipment, while also addressing public perceptions and concerns through publicity campaigns
and demonstration programs. This chapter presents technical information surrounding the physical properties of hydrogen and its
characteristic behavior during leakage, combustion, and explosion. It also details current international standards in this area
surrounding the safe use of hydrogen as a fuel.

4.01.2.5

Chapter 4.05: Hydrogen Storage: Compressed Gas

A critical element in the development of a potential ‘hydrogen economy’ is the ability to store and transport hydrogen gas as a fuel.
The storage of hydrogen has, and remains most commonly, been in the form of a compressed gas. Historically, this has typically
been at ‘industrial’ gas pressures of around 200 bar, a typical pressure for a steel cylinder of hydrogen for laboratory or factory use;
however, in recent times, the use of high pressure (up to 700 bar) storage of hydrogen within composite pressure vessels has been
promoted particularly for vehicular use. This chapter explores the issues and technologies involved in storing hydrogen gas at such
pressures, and discusses the safe handling and use of hydrogen within such systems.

4.01.2.6

Chapter 4.06: Hydrogen Storage: Liquid and Chemical


Hydrogen storage in the form of compressed gas has a number of limitations and this chapter studies other possible forms of storing
hydrogen as a fuel, namely, as a liquid at very low temperatures or within chemical media capable of absorption or reversible
material reactions. Some details of these alternative forms of storage are presented, and the issues surrounding reversibility, the
dynamics and speed of storage/release, cost, storage density and cyclability are discussed within this context.

4.01.2.7

Chapter 4.07: Alkaline Fuel Cells: Theory and Application

There are several different types of fuel cell that can be broadly characterized by the type of electrolyte used within the cell. In the late
1950s, the main electrolyte of interest was potassium hydroxide, an alkaline solution, and hence, it was this particular fuel cell technology
that reinvigorated research and commercial interest following on from successful use in early space flights. This chapter presents both
further technical detail of alkaline fuel cell (AFC) technology and further historical information on past and current developments.

4.01.2.8

Chapter 4.08: PEM Fuel Cells: Applications

For transport applications of fuel cells, an obvious disadvantage to the use of AFCs was the requirement to contain and protect (in the
event of collision/accident) a highly caustic liquid electrolyte. A solution to this issue arose from the study of ion-conducting solid
membranes, which obviated the need for a liquid electrolyte, thereby removing accompanying spill/leakage issues. This work led to the
development of positive ion-conducting membranes (e.g., proton conduction), as opposed to negative hydroxyl ion conduction within
AFC, and spawned the PEM fuel cell sector. This genre of fuel cell is now the most widely researched and most promising for utilization
within the transport sector, and this chapter explores the material developments, capabilities, and demonstration of this technology.

4.01.2.9

Chapter 4.09: Molten Carbonate Fuel Cells: Theory and Application


As the chapter name suggests, this fuel cell technology is characterized by operation using a molten carbonate salt as the cell
electrolyte. This requires temperatures of several hundred degrees Celsius to produce the molten electrolyte and achieve appropriate


8

Preface and Context to Hydrogen and Fuel Cells

ion mobility, thereby enabling fuel cell operation, and this technology has developed to accommodate the resulting technical
challenges. This chapter presents details of the development of molten carbonate fuel cells, in particular the significant installed
commercial capacity of these fuel cells, and discusses the future potential for this technology.

4.01.2.10

Chapter 4.10: Solid Oxide Fuel Cells: Theory and Materials

Of the range of electrolytes available for use within fuel cells, the highest temperature electrolyte currently used is a negative
ion-conducting ceramic within the solid oxide fuel cell (SOFC) class. Typically, operating temperatures of between 550 and 1000 °C
are needed to produce the conditions necessary for appropriate ionic conduction in these materials; however, such high-temperature
fuel cell technology is attractive in many process industries where high-grade waste heat may be employed to readily create the operating
temperatures or direct use of fossil fuel combustion products that are available. This chapter explains how an SOFC operates, discusses
the material requirements and developments, and presents information on the current state of the art in this sector.

4.01.2.11

Chapter 4.11: Biological and Microbial Fuel Cells

Microbial and biological fuel cells are relatively new forms of fuel cell that typically use either hydrocarbons as a fuel, or utilize
electrogenic bacteria to convert chemical energy to electrical energy rather than a more typical electrocatalyst. These forms of fuel cell
are gaining prominence in developing niche markets, such as wastewater treatment where electrical energy can be derived while

simultaneously processing and cleaning the waste stream, and are potentially able to address the growing legislative requirements to
clean hitherto neglected process streams. This chapter explores the theory and materials (including enzymes and bacteria) employed
to create these unique fuel cells and presents details of their resulting performance characteristics.

4.01.2.12

Chapter 4.12: Hydrogen and Fuel Cells in Transport

Of the many application areas for fuel cells, the sector that may provide the greatest advance in terms of commercial breakthrough and
success is within the transport field. Fuel cell utilization within electric vehicles, ships or as on-board electrical generators for load
refrigeration or cab power, is increasing rapidly and offers efficiency and carbon emissions benefits compared to conventional fossil
fuel options. The transport sector offers both a huge market and a direct interface to the public (compared to the stationary power
market for fuel cells which tends to ‘isolate’ the consumer from the technology, whereas, for example, transport fuel cells are accessible
under the hood of a car); hence, fuel cells can gain rapid consumer acceptance and market traction within this key end-use sector. This
chapter presents details of fuel cell technology relevant to the transport sector and illustrates several examples of typical use.

4.01.2.13

Chapter 4.13: H2 and Fuel Cells as Controlled Renewables: FC Power Electronics

Neither hydrogen nor fuel cells are formally defined as ‘renewable sources’ of power within current legislation [7]; however, they are
commonly viewed as an integral element of any future clean energy mix. Uniquely, within the context of renewable energy sources
where many of these sources are stochastic (i.e., variable) and uncontrolled in terms of their energy delivery, hydrogen and fuel cell
technology offers the user direct control of the power delivered at any instant. The ability to control the power and energy delivery
from a fuel cell, and to control the storage and utilization of hydrogen as a fuel, conveys significant benefits to the utilization of
hydrogen and fuel cell technologies. This chapter explores the use of power electronics as a mechanism to control the output
electrical power from a fuel cell, and examines the type and nature of the power electronic interface between the electrochemical cell
and the electrical load.

4.01.2.14


Chapter 4.14: Future Perspective on Hydrogen and Fuel Cells

The final chapter of this volume presents a perspective of how hydrogen and fuel cell technology may develop in the future, and
how hydrogen use as a fuel may pervade more aspects of our lives. This concept, of a future hydrogen-based economy, is a view that
many researchers and analysts believe it is inevitable as the world’s fossil fuel resources are depleted and environmental pressures
compound the shift toward a cleaner energy supply.

4.01.3 Hydrogen and Fuel Cell Technology – Supplementary Material
It has not been possible within this volume to adequately cover ‘all’ aspects of hydrogen and fuel cell technology, and this brief section aims
to introduce the reader to areas considered by the author as important with appropriate references to allow further reading as required.

4.01.3.1

Flow Cells or Regenerative Fuel Cells

The first technology of interest is the flow cell or flow battery, where the energy capacity of the system is stored external to the cell
generally in the form of oppositely charged liquid electrolytes. These electrolytes are reacted within a fuel cell, in the form of two half


Preface and Context to Hydrogen and Fuel Cells

9

cell reactions (see Section 4.01.4.1) separated by an ion-selective membrane (i.e., a membrane that allows only specific ions to pass
through it but prevents the liquid electrolytes from mixing directly) that convert the chemical energy in the electrolytes fuels to
electrical energy. This technology is viewed as different from a ‘conventional’ hydrogen and oxygen (or air) fuel cell. This technology
can be encountered under a number of different names, for example, regenerative fuel cell or redox flow fuel cell.
This technology is exemplified by the technology such as the sodium polysulfide bromide-based system developed by a UK
company called Regenesys, from the late 1990s to the early 2000s. Regenesys aimed at commissioning a demonstration plant rated

at 120 MWh, 12 MW at a site near Little Barford, Cambridgeshire, UK, to evaluate regenerative fuel cell technology as a large-scale
form of electrical energy storage. The test site was almost completed, as per Figure 8, but ultimately never fully commissioned, as
explained in Reference [8].
Figure 9 shows the actual regenerative fuel cell modules within the Little Barford test site. As with other forms of fuel cell, this
technology produces DC electrical power, and hence, a power electronic inverter was required to convert this to AC electrical power
suitable for delivery into the electricity grid. For the Little Barford test site, the inverter was developed by ABB and was rated at
18.25 MVA [9] although the fuel cells were rated to deliver only up to 12 MW.
The Regenesys technology employed two electrolytes: sodium bromide (NaBr) and sodium polysulfide (NaS2). These two liquid
electrolytes were circulated through the regenerative fuel cell as per the diagram shown in Figure 10, where the half-cell redox
reactions occurred, generating electrical energy if pumped in one direction, and capable of reversal to store electrical energy if the
electrolytes were pumped through the cell in the opposite sense (with electrical energy input to the cell in this instance).

Figure 8 Exterior view of Regenesys Technologies Ltd. pilot utility scale energy storage plant at Little Barford, Cambridgeshire, UK, July 2003. From
“Regenesys Utility Scale Energy Storage – Project Summary”, DTI, Contract Number: K/EL/00246/00/00, URN Number: 04/0148, 2004.

Figure 9 Interior view of the plant – stream of XL modules. From “Regenesys Utility Scale Energy Storage – Project Summary”, DTI, Contract Number: K/
EL/00246/00/00, URN Number: 04/0148, 2004.


10

Preface and Context to Hydrogen and Fuel Cells

Electrolyte
tank
Electrode

Ion-selective
membrane



Regenerative

fuel cell


Electrolyte

Pump

Power source/load


Electrolyte
tank

Electrolyte


Pump


Figure 10 Flow diagram of Regenesys flow cell. Price, A, et al. A novel approach to utility scale energy storage. IEE Power Engineering Journal,
June 1999.

This technology, along with other chemical types such as the zinc bromide battery [10], vanadium redox battery [11], and iron–
chromium [12] system, can all act as types of redox flow cell, and the reader is directed to the various websites for further information.

4.01.3.2


Hydrogen Production – Electrolysis

Currently, the most common method of producing hydrogen gas on a large scale is through the process of steam methane
reformation, where methane gas (CH4) is reacted with steam (H2O) to, ultimately, produce carbon dioxide (CO2) and hydrogen
(H2). However, within the context of this volume, focussing on renewable energy, the preferred method of producing hydrogen is
via electrolysis of water (see Section 4.01.1) where electrical energy, preferably from clean sources such as wind or solar power, is
used to split water into its constituent gases, namely, hydrogen and oxygen.
This process, and accompanying electrolyser technology, has been extensively developed by companies such as Norsk Hydro
[13] which, since 1920s, has been using cheap hydroelectric electricity in Norway to split water into hydrogen and oxygen gases, and
then combine the hydrogen with nitrogen extracted from the air, to form ammonium, a key component of fertilizer.
The technology surrounding large-scale electrolysers, even today, is based on alkaline technology employing potassium hydro­
xide as an electrolyte, and such units are available in capacities up to 2 MW rating, as shown in Figure 11. Figure 11 shows the
electrical connections on the left-hand side of this image to supply DC electrical power to the fuel cells, with the white lye
(potassium hydroxide) tanks shown toward the rear right-hand side of the image.

Figure 11 NEL Hydrogen, 2 MW, alkaline electrolyser. NEL Hydrogen Ltd., www.nel-hydrogen.com (accessed March 2012) [14].


Preface and Context to Hydrogen and Fuel Cells

Table 1

11

Specification of NEL Hydrogen electrolysers

Capacity
Capacity range (Nm3 H2h−1)
Maximum Nm3 H2 per cell
Energy

Power consumption at 4000 amp DC (kWhNm−3 H2)
Power consumption at 5150 amp DC (kWhNm−3 H2)
Purity
H2 purity (%)
O2 purity (%)
H2 purity after purification (%)
Pressure
H2 outlet pressure after electrolyser
Maximum H2 outlet pressure after compressor
Operation
Operating temperature
Operation
Electrolyte
Feed water consumption

10–485
2.11
4.1 Æ 0.1
4.3 Æ 0.1
99.9 Æ 0.1
99.5 Æ 0.1
99.999 8% (2 ppm)
200−500 mm WG
440 bar g
80 °C
Automatic, 20–100% of max capacity
25% KOH aqueous solution
0.9 l(Nm−3H2)−1

Source: NEL Hydrogen Ltd.


The technical specification of such a large-scale alkaline electrolyser is shown in Table 1, which defines the energy requirement
for hydrogen production (e.g., 4.1 kWh Nm−3 of H2) and the feed water flow rate required (e.g., 0.9 l Nm−3 of H2). It also defines the
typical purity of hydrogen produced and the typical use of subsequent cleanup processes to improve the hydrogen gas purity to five
‘9s’ purity, that is, 99.999% which is typically required for fuel use for fuel cells.
The flow diagram for this technology is shown in Figure 12, including the electrical supply, electrolyte, and gas separation and
handling systems.
Figure 13 shows a typical electrode assembly for a large-scale electrolyser (e.g., for a unit such as that shown in Figure 11). This
image also shows the manifold ducts on the top and bottom of the electrode which are used to capture the gases and circulate the
liquid electrolyte, respectively.
Figure 14 illustrates the typical separator membrane that is used which permits ionic conduction between the anode and
cathode electrodes (see Section 4.01.4.1) yet prevents gas and liquid crossover. Figure 15 gives a closeup view of the gas manifold
ducts, illustrating (from an edge view) that the oxygen and hydrogen gases are produced and collected from opposite sides of the
membrane and fed to separate gas ducts.
Figure 16 shows an image of the electrolyser units within the Glomfjord ammonia plant in Norway, as a series of individual
alkaline electrolysers of a type similar to those in Figure 11.

4.01.3.3

Hydrogen Demonstration Units – State of the Art

There are a number of hydrogen production and utilization demonstration units in trial around the world, based on a variety of
different fuel cell technologies. One such system, which has attracted a lot of attention within the United Kingdom, is the ITM Power

Oxygen to atmosphere


Water
Oxygen


Hydrogen

Electrolyte


Deoxo

To process

High voltage supply

−

+

O2

H2

High pressure
storage

Feed water
Dryer

Transformer Rectifier
Electrolyser

Gas/Lye
separator


Gas/Lye
separator

Scrubber

Feed water to
electrolyser
Lye tank

Figure 12 Flow diagram of alkaline electrolyser (NEL Hydrogen Ltd.).

Gas holder Compressor


12

Preface and Context to Hydrogen and Fuel Cells

Figure 13 Electrode of alkaline electrolyser (NEL Hydrogen Ltd.).

Figure 14 Non-asbestos diaphragm for alkaline electrolyser (NEL Hydrogen Ltd.).

Hydrogen

Oxygen
Figure 15 Gas ducts on alkaline electrolyser (edge view of cell, with ducts on either side of diaphragm) (NEL Hydrogen Ltd.).


Preface and Context to Hydrogen and Fuel Cells


13

Figure 16 Hydro’s electrolyser plant for ammonia-production in Glomfjord, Norway (NEL Hydrogen Ltd.).

Figure 17 Water treatment. Courtesy of Dr. A. Cruden.

[15] HOST (Hydrogen On-Site Trial) unit, which demonstrates on-site hydrogen production via PEM electrolyser technology,
compression and high pressure storage of hydrogen gas, and finally dispensing of this fuel gas to on-vehicle storage for use within a
hydrogen internal combustion engine (HICE) on a van.
Some of the technology used within the HOST unit is illustrated in the following series of images. For example, Figure 17 shows
the water processing plant requiring to clean up a typical potable water supply for use within high purity PEM electrolysers.
Figure 18 shows a palladium dryer unit that is used to help purify the hydrogen gas produced from the PEM electrolysers prior to
gas compression and storage. This stage helps prevent corrosion issues within the compressor and storage vessels.
The purified hydrogen gas is subsequently stored in a tiered system of pressure vessels from ‘low’ pressure (up to 250 bar) to
‘medium’ pressure (up to 350 bar, shown in Figure 19) tanks, and finally up to 450 bar in the ‘high’ pressure tanks (the top four
tanks shown in Figure 20). The compressor used is shown in Figure 21.
The high pressure hydrogen stored is subsequently dispensed via a high pressure nozzle connector, as shown in Figure 22, which
directly mates to a vehicle-mounted receptacle, shown in Figure 23. A vehicle refueller control system (i.e., a dispensing refueller
system) is shown in Figure 24, and is used to control and monitor the quantity of hydrogen fuel gas dispensed to the vehicle.
Figure 24 effectively illustrates a future hydrogen refuelling pump, akin to a conventional ‘petrol pump’ that would be found on any
petrol station forecourt. Hence, Figure 24 shows the type of technology that is being developed for use within a future hydrogen
economy for dispensing vehicle fuel.


14

Preface and Context to Hydrogen and Fuel Cells

Figure 18 Dryer/heater/separator. Courtesy of Dr. A. Cruden.


Figure 19 Storage tanks. Courtesy of Dr. A. Cruden.

Figure 25 illustrates a typical commercial van that employs a modified spark ignition combustion engine that can utilize
hydrogen gas as a fuel. Thus, the HOST trial system demonstrates and evidences all the technology required to produce, compress,
store, dispense, and utilize hydrogen as a clean vehicle fuel. Such systems are being trialled around the world [16] and highlight the
progress that has been made in the underlying technology areas.


Preface and Context to Hydrogen and Fuel Cells

15

Figure 20 High pressure storage cylinders and bottom buffer store. Courtesy of Dr. A. Cruden.

Figure 21 Compressor. Courtesy of Dr. A. Cruden.

A further image of a test site is the NEL Hydrogen park at Porsgrunn (see Figure 26) in Norway, where demonstration
of latest generation alkaline electrolyser technology, compression and storage, and vehicle refuelling equipment is under
active trial.


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Preface and Context to Hydrogen and Fuel Cells

Figure 22 WEH refueller nozzle. Courtesy of Dr. A. Cruden.

Figure 23 Compressed hydrogen gas vehicle receptacle. Courtesy of Dr. A. Cruden.



Preface and Context to Hydrogen and Fuel Cells

17

Figure 24 Hydrogen compressed gas vehicle refuelling station. Courtesy of Dr. A. Cruden.

Figure 25 Hydrogen internal combustion engine (HICE) van. Courtesy of Dr. A. Cruden.

4.01.4 Introduction to Basic Electrochemistry
As discussed earlier, a hydrogen fuel cell is effectively a ‘continuously operating battery’; that is, as long as ‘fuel gases’ (in this case,
hydrogen and oxygen) are fed to the anode and the cathode, respectively, the fuel cell will continue to produce DC electrical power
at its terminals. In particular, a fuel cell performs a direct energy conversion from the chemical energy in the fuel gases to electrical
energy. This differs from conventional electricity generation within a thermal power station, say coal or gas fired, where the chemical
energy in the fuel is first converted to heat via combustion, and the heat is used to produce superheated high pressure steam within a
boiler, which in turn is then used to produce rotating kinetic energy via a steam turbine, which in turn then drives the main rotating
shaft of a synchronous AC electrical generator. As there are several energy conversion stages in the conventional thermal power
station process, each with their own loss mechanism, electricity produced by this process is typically ∼40% efficient, that is, typically


18

Preface and Context to Hydrogen and Fuel Cells

Figure 26 NEL Hydrogen Porsgrunn Hydrogen Station. Courtesy of Dr. A. Cruden.

60% of the input energy in the primary fuel (coal or gas) is ultimately not converted to electrical energy, but is lost, mostly in the
form of waste heat. Finally, this rather poor efficiency can only be achieved by building thermal power stations at a large scale,
typically >1000 MW rating, as the thermal losses tend to increase at a smaller scale.
As a fuel cell undertakes a direct energy conversion, its efficiency is higher, typically in the range of ∼50% for conversion of

chemical to electrical energy. Further, this efficiency is broadly independent of the scale of the fuel cell, from a few hundred watts to
several megawatts, leading to a highly efficient modular electricity generating technology.
The basic electrochemical theory of a fuel cell will now be introduced.

4.01.4.1

Redox Reactions

To begin with some electrochemical terms requires definition. First, an ‘oxidation’ process is defined as the addition of oxygen or
removal of hydrogen from a substance. A ‘reduction’ process is the addition of hydrogen or removal of oxygen from a substance,
that is, the opposite of the oxidation process. Further, no oxidation process can take place without a corresponding reduction
reaction. For example,
Oxidation

½1Š

2H2 + O2 ⇒ 2H2O
Reduction

In eqn [1], the hydrogen is oxidized, that is, oxygen is added to form water, while simultaneously the oxygen is reduced, that is,
hydrogen is added to form water. On this basis, the oxygen acts as the ‘oxidizing agent’, while the hydrogen acts as the ‘reducing agent’.
However, there is a more general form of redox reaction which can be further explained by way of another example reaction (see
eqn [2]):
Oxidation

H2S + CI2 ⇒ 2HCI + S

½2Š

Reduction


In eqn [2], it is clear that the hydrogen sulfide is the oxidizing agent, that is, it removes hydrogen to form sulfur; however, according
to the premise stated above, there must be an accompanying reduction reaction. Hence, the chlorine must act as the reducing agent.
An alternative definition of redox reactions may be formed from eqn [2] if the ionic reaction equations are considered:
H2 S þ Cl2 ⇒2Hþ Cl− þ S
H2 S ⇒2Hþ þ S þ 2e−
Cl2 þ 2e− ⇒2Cl−

oxidation
reduction

½3Š


Preface and Context to Hydrogen and Fuel Cells

19

A

KNO3/H2O

Cu

Zn

CuSO4/H2O

ZnSO4/H2O


Figure 27 Simple zinc/copper electrochemical cell. />
From the ionic reactions, it is clear that the oxidation reaction occurs with both the removal of hydrogen from the hydrogen sulfide
and the ‘loss of electrons’. Similarly, the reduction reaction occurs with the ‘addition of electrons’.
This therefore leads to a more general definition of redox reactions and following mnemonic to remember this:
Oxidation is the removal of electrons from a substance.

Reduction is the addition of electrons to a substance.

‘LEO says GER – Lose Electrons Oxidation, Gain Electrons Reduction’

All such ionic reactions can be realized under conditions where an electric current, that is, the flow of electrons, is present in an
external electric circuit (as opposed to an internal current within a closed beaker say). Under these conditions, an ‘electrochemical
cell’ is created. A simple realization of this is a basic zinc/copper electrochemical cell, formed with a zinc rod placed in a beaker of
solution of its own ions (say zinc sulfate, ZnSO4), and similarly a copper rod is placed in a beaker of its own ions (say copper sulfate,
CuSO4). Connecting the two beakers is a ‘salt bridge’, a device used to allow ionic (as opposed to electronic) conduction between
the two ‘half cells’. This arrangement is shown in Figure 27.
From Figure 27, it is found that the ammeter will register current flow, the flow of electrons, from the zinc electrode to the copper
electrode. This current flow occurs while the zinc electrode slowly dissolves into the solution, and the copper electrode undergoes
plating of fresh copper deposits on its surface. The corresponding redox ionic reactions are as follows (also know as the ‘half-cell
reactions’):
Zn þ Cu2 þ ⇒Zn2 þ þ Cu
Zn ⇒Zn2 þ þ 2e−
oxidation − anode
Cu2 þ þ 2e− ⇒Cu
reduction − cathode

½4Š

Equation [4] indicates that the zinc is being oxidized, while the copper is being reduced, and the electrons are ‘lost’ from the zinc
flow through the external electrical circuit to the copper electrode where they are ‘added’ to the copper ion as part of the reduction

process. Within the fields of electrochemistry and physics, convention states that the electrode where oxidation occurs is termed the
‘anode’ and the electrode where reduction occurs is termed the ‘cathode’. Hence, for a galvanic cell (i.e., battery or fuel cell), the
anode is the negatively charged electrode (from Figure 27 and eqn [4], the zinc electrode ‘dissolves’ to a solution of zinc ions (Zn2+)
and becomes negatively charged due to accumulation of electrons), and the cathode is the positively charged electrode (as the
copper ions (Cu2+) in solution accept electrons and deposit copper on the electrode surface, creating a deficit of electrons (a positive
charge)).
The amount of both zinc lost and copper plating added is proportional to the amount of electrical charge (i.e., electrons) that
flows between the electrodes, where the amount of electrical charge, Q (coulombs), is given by
Q ¼ IÃ t

½5Š

where I is the current (amps) and t is the time (seconds).
If the electrical current is not allowed to flow (e.g., by breaking the electrical circuit via the ammeter), then there is no loss of zinc
nor fresh plating of the copper electrode, proving that the electron transfer of the equations in [4] is intrinsic to the redox process. If


20

Preface and Context to Hydrogen and Fuel Cells

+

3V

1

−

2


3
S

R

Resistance wire
A

A

A

Ammeter

−

+

E-cell
Sliding contact
Cu/Cu2+⏐⏐Zn2+/Zn
Figure 28 Experimental setup to determine the open-circuit potential of an E-cell.

the ammeter in Figure 27 was replaced by a voltmeter, a potential difference (voltage) of 1.1 V would be measured between the two
electrodes.
An alternative experimental arrangement that helps illustrate the operation of an electrochemical cell (E-cell), and determine the
potential difference of the cell, is shown in Figure 28.
The 3 V battery shown in Figure 28 drives current through the resistance wire (RS). One terminal of the E-cell is connected to the
resistance wire (R) and the other terminal of the E-cell is connected to the resistance wire via a sliding contact. By varying the

position of the sliding contact, a variable potential (voltage) is applied to the E-cell terminals. At position ‘2’, it is assumed that the
applied voltage via the battery and resistance wire is exactly equal and opposite to the potential difference generated by the E-cell, at
which point the ammeter shows zero current. At this position, measuring the voltage across the E-cell (voltage across R2) is a direct
measure of the ‘open-circuit’ potential of the cell. This is characterized as the potential where zero electronic current flows. For the
simple zinc/copper E-cell of Figure 27, this equals a voltage of 1.1 V.
Considering the equations in [4], electrochemically this means that the reactions are occurring at exactly the equal and opposite
rate, for example, for each molecule of zinc that is oxidized to Zn2+ and liberates two electrons, a molecule of zinc ions (Zn2+)
combines with two electrons (2e−) to produce zinc (Zn). The arrows in eqns [6] and [7] indicate these ‘reversible’ reactions, and at
open circuit, these competing reactions are balanced, that is, they are at ‘equilibrium’.
Zn ⇔Zn2 þ þ 2e−

½6Š

Cu2 þ þ 2e− ⇔Cu

½7Š

If the sliding contact is now moved to position ‘1’, the ammeter will now show current flow from the E-cell, and the voltage
measured across the E-cell terminals will be <1.1 V. At this position, the reaction in eqn [6] will progress much faster in the forward
(right hand) direction, thereby liberating electrons which will flow from the zinc side of the cell, via resistance wire R1 to the copper
side of the cell, as shown in eqn [8]. This indicates the cell is no longer at equilibrium.
Zn ⇒Zn2 þ þ 2e−

½8Š

At the same time, the copper ions in solution will accept these electrons and deposit fresh copper plate on the copper electrode
surface, according to eqn [9]:
Cu2 þ þ 2e− ⇒Cu

½9Š


Under the conditions of eqns [8] and [9], the E-cell is producing electric current and is acting like a battery; that is, the E-cell causes a
flow of electrons from the cell and is converting the chemical energy in the zinc and copper to electrical energy.
Conversely, if the sliding contact is moved to position ‘3’, the ammeter will show current flow to the E-cell (from the 3 V battery)
and the voltage measured across the E-cell terminals will be >1.1 V. At this position, the reaction in eqn [6] will progress much faster
in the reverse (left-hand) direction, and zinc ions from the solution will combine with free electrons from the supplied electric
current to deposit fresh zinc on the electrode surface. This is shown in eqn [10]:
Zn2 þ þ 2e− ⇒Zn

½10Š

Similarly, the copper electrode reaction in eqn [7] will also be driven in the reverse (left-hand) direction, and the copper electrode
will gradually dissolve into its salt solution, as shown in eqn [11]:
Cu ⇒Cu2 þ þ 2e−

½11Š


Preface and Context to Hydrogen and Fuel Cells

4.01.4.2

21

Electrochemical Series

The voltage measured between the electrodes of the zinc/copper cell, shown in Figure 27, on open circuit was measured as 1.1 V. If
both the electrode materials were changed, then it is highly likely that a different voltage would be measured, although it would not
be easy to readily compare the performance of these two electrochemical cells as there is no common reference between them.
Indeed, the need to allow comparative measure of the electrochemical performance of different materials led to the development of

the ‘standard hydrogen electrode’ (SHE) that is given an arbitrary potential of 0.00 V (Figure 29).
The SHE is given the potential of 0 V only at certain conditions: a hydrogen gas pressure of 1 bar, bubbling over a platinum
electrode foil, immersed in a solution of 1 molar H+ ions (i.e., acid), at a temperature of 25 °C (298 K). This is embodied in eqn [12]:
−
2Hþ
ðaqÞ þ 2e ⇔H2 ðg Þ

½12Š

0:00 V

By definition of this arbitrary reference, all other materials can be electrochemically compared to the SHE, for example, magnesium,
as shown in Figure 30. The voltmeter in Figure 30 will measure an open-circuit potential of 2.37 V, with the magnesium electrode

Hydrogen
at 1 bar

Temperature = 298 K

Platinum wire

Platinum foil
covered in
porous platinum
Dilute sulfuric acid
[H+] = 1 mol dm−3
Figure 29 Example of a standard hydrogen electrode. />
V
Temperature = 298 K
High resistance

voltmeter

Hydrogen
at 1 bar

Magnesium

Platinum wire

Salt
bridge

Platinum foil

covered in

porous platinum

Dilute sulfuric acid
[H+] = 1 mol dm−3

Magnesium sulfate
solution
[Mg2+] = 1 mol dm−3

Figure 30 Use of an SHE within an electrochemical cell to determine the potential of magnesium. />introduction.html


22


Preface and Context to Hydrogen and Fuel Cells

Table 2

Series of electrochemical potentials

E°
(V)

Reaction
Li ⇔ Li+ + e−
Mg ⇔

Mg2+

−3.04
−

+ 2e

Al ⇔ Al3+ + 3e−
2H2 + 4OH− ⇔ 4H2O + 4e−

−2.37
−1.66
−0.83

Zn ⇔ Zn2+ + 2e−

−0.76


Fe ⇔ Fe2+ + 2e−

−0.44

−

−0.25

+ 2e

Ni ⇔

Ni2+

H2 ⇔

2H+ + 2e−

Cu2+ + 2e− ⇔ Cu
Ag+ + e− ⇔ Ag

Standard hydrogen electrode

+0.40
+0.80

O2 + 4H+ + 4e− ⇔ 2H2O
⇔ Au


0.00
+0.34

O2 + 2H2O + 4e− ⇔ 4OH−

Au3+ + 3e−

Most likely to oxidize

+1.23
Most likely to reduce

+1.50

being determined (as per the experiments using the apparatus in Figure 27) as the electrode experiencing oxidation, that is, the
magnesium, liberates electrons and is the anode, as per eqn [13]:
Mg ⇔Mg2 þ þ 2e−

−2:37 V

½13Š

If the electrochemical cell of Figure 30 is used as a battery, then the overall electrode reactions are as shown in eqn [14]:
Mg ⇒Mg2 þ þ 2e− −2:37 V
−
0:00 V
2Hþ
ðaqÞ þ 2e ⇒H2 ðg Þ

½14Š


The testing of a range of different electrode materials has been undertaken and defined in Tables of Electrochemical Potentials, see Table 2.
Table 2 is shown with all reactions at open circuit; however, the preferred direction of each reaction versus a SHE is as indicated,
reading each equation from left to right. Further, the reactions at the top of the table indicate materials best suited for oxidation, that is, the
strongest oxidizing agents, while materials at the bottom of the table are best suited for reduction, that is, the strongest reducing agents.
This table can also be used to determine the open circuit potential, and anode and cathode, of any given electrochemical cell. For
example, taking zinc and copper again (as per Figure 27), Table 2 indicates that zinc will oxidize (i.e., form the anode of the cell)
with a standard potential of –0.76 V, while copper will reduce and form the cathode at a standard potential of +0.34 V. Hence,
overall, the zinc/copper electrochemical cell will produce an open-circuit potential of +0.34 – (–0.76) = 1.1 V (as noted previously).
Similarly, taking an aluminum/zinc cell, this time the aluminum will oxidize and form the anode at a standard potential of
–1.66 V, while this time the zinc will reduce and form the cathode at a standard potential of –0.76 V. Hence, overall, the aluminum/
zinc electrochemical cell will produce an open-circuit potential of –0.76 – (–1.66) = 0.9 V.

4.01.4.3

Gibbs Energy – Useful Work

The equation for power, P (watts), produced by an electric circuit working with a potential difference of V (volts) and current I
(amps) is given by eqn [15]:
P ¼ V ÃI

½15Š

The energy E (joules) consumed while the electric circuit is working at this power is given by eqn [16]:
E ¼ PÃ t

½16Š

where t is the time (seconds) that the circuit is working for. Hence, by substituting [15] in [16], and using eqn [5]
E ¼ P Ã t ¼ V Ã IÃ t ¼ V Ã Q


½17Š

Hence, the units of energy (joules, J) can be expressed in terms of volts. Coulombs (VC), or 1 joule of energy, is required to move 1
coulomb of charge through a potential difference of 1 volt.
Hence, as the open circuit voltage for the zinc/copper E-cell is 1.1 V, and from eqns [6] and [7], it is clear that there are two
electrons involved in the accompanying redox reactions, then for 1 mole of zinc and 1 mole of copper reacting to completion in an
E-cell as per Figure XX, would result in a useful energy production per mole given by


Preface and Context to Hydrogen and Fuel Cells
Wuseful ¼ nFE
¼2Ã 96485Ã 1:1 ¼ 212267 Joules ¼ 212:3 kJ mol − 1

23

½18Š

where n is the number of moles of electrons that flow during the reaction (= 2 from eqns [6] and [7]), F is the Faraday constant that
defines the total electric charge per mole of a substance (96 485 °C mol−1), and E is the open-circuit voltage of the E-cell (V).
Equation [18] is of the same form as eqn [17], where nF is equivalent to Q, the total charge transferred.
From classic thermodynamics, the useful work, Wuseful, is more commonly denoted as ΔG, the ‘Gibbs free energy’ where, by
convention, a decrease in free energy, –ΔG, occurs when useful work is done (i.e., the zinc/copper E-cell reaction causes an electric
current to flow that can be used for performing ‘useful work’, e.g., powering a light bulb); hence, eqn [18] can more correctly be
rewritten as
−ΔG ¼ Wuseful ¼ 212:3 kJ mol − 1
ΔG ¼ −212:3 kJ mol − 1

½19Š


As ΔG is negative, by definition, useful work has been carried out, and the reaction will tend to occur in this sense. By contrast, if a
chemical reaction results in a positive ΔG, then this implies that the reaction is not feasible without input of energy, and it is not
feasible to produce ‘useful work’ via such a reaction.
The first law of thermodynamics states that energy must be conserved, that is, that energy can change state (i.e., from chemical
energy to heat) but cannot be created nor destroyed. As the internal energy of a system cannot be determined in an absolute sense,
then the first law states that the change in the internal energy, dU, of a system can be expressed as the change in heat, dQ, and the
change in useful work, dW.
Mathematically, this can be expressed as
dW ¼ dU− dQ

½20Š

For a reversible process in a closed system, eqn [20] can be rewritten in terms of the change in the total energy of the system, more
commonly known as the change in ‘enthalpy’ of the system, ΔH (kJ mol−1); the change in useful work also known as the Gibbs free
energy, ΔG (kJ mol−1); and the change in heat of the system calculated as dQ = TΔS, where T is the temperature in units of kelvin (K)
and ΔS is the change in ‘entropy’ of the system (J (K mol)−1) [17]:
ΔG ¼ ΔH −TΔS

½21Š

Each system, thermodynamically, is trying to achieve the minimum of free energy, that is, it is trying to achieve its lowest, most
stable energy state. For example, eqn [21] tells us that reactions that release free energy (i.e., ΔG is a ‘negative’ value) are generally
spontaneous and result in more stable resulting state, whereas reactions that result in a ‘positive’ ΔG require energy to be input to the
system to force the reaction to occur and generally result in system that is less stable.
Values for the enthalpies and entropies of reactants and products for a given reaction can be found in many data tables, and the
calculation of ΔH and ΔS can be best illustrated by way of an example, considering the zinc/copper E-cell from Section 4.01.3.1. The
reaction for the zinc/copper E-cell is
Zn ðs Þ þ CuSO4 ðaq Þ ⇒ ZnSO4 ðaq Þ þ Cu ðs Þ

½22Š


The values of enthalpy and entropy for these materials and compounds are [18] as follows:
SZn ðs Þ ¼ 41:6 J ðK mol Þ − 1
SZnSO4 ðaq Þ ¼ 120 J ðK mol Þ − 1
SCu ðs Þ ¼ 33:2 J ðK mol Þ − 1
SCuSO4 ðaq Þ ¼ 109:2 J ðK mol Þ − 1
ΔHZnSO4 ¼ −982:8 kJ mol − 1
ΔHCuSO4 ¼ −771:4 kJ mol − 1
Hence, the values for ΔS, TΔS, and ΔH, at a temperature of 25 °C (298 K) are as follows:
�
� �
�
ΔS ¼ Sproducts –Sreactants ¼ SZnSO4 ðaq Þ þ SCu ðs Þ – SZn ðs Þ þ SCuSO4 ðaq Þ
¼ð120 þ 33:2Þ–ð41:6 þ 109:2Þ ¼ 153:2–150:8 ¼ 2:6 J ðK mol Þ − 1
TΔS ¼ 298Ã ð2:6Þ ¼ 774:8 J mol − 1
ΔH ¼ ΔDHZnSO4 − ΔHCuSO4 ¼ −982:8–ð−771:4Þ ¼ −211:4 kJ mol − 1
Replacing these values for TΔS and ΔH in eqn [21] gives
ΔG ¼ ΔH− TΔS ¼ −211400 –774:8 ¼ −212174:8 ¼ −212:2 kJ mol − 1

½23Š

This is the same as the value for ΔG (within the accuracy of the thermodynamic data from the data tables) from eqn [19], however,
derived from fundamental properties. In this case, the value of dQ = TΔS is positive, implying that a small amount of heat is actually


Preface and Context to Hydrogen and Fuel Cells

Electrolyte
aqueous
solution

K+

OH−

O2

H2O

O2 compartment

+

Electrode

Electrode

−

H2

H2 compartment

H2O

24

2H2 + 4OH − ⇒ 4H 2O + 4e−

O2 + 2H2O + 4e− ⇒ 4OH −
Figure 31 Operation of an AFC.


absorbed from the surroundings during this reaction and that a significant amount of energy is available as useful work (in this case,
electrical energy).
However, consider the theoretical case of the alkaline hydrogen fuel cell, as first illustrated in Figures 6 and 7. The AFC works
commonly employing a potassium hydroxide electrolyte, where the negative hydroxide ion (OH−) enables internal fuel cell
operation by facilitating charge transfer between cathode and anode as illustrated in Figure 31.
With reference to eqn [24] and Figure 31, it is evident that the hydrogen gas, delivered to the H2 compartment, permeates through
the electrode membrane assembly and reacts with the hydroxyl ion, producing water (which exits the cell from the H2 compartment)
and releases electrons. Hence, the hydrogen electrode becomes the negative electrode, or anode, where oxidation occurs.
Further, with reference to eqn [25] and Figure 31, it is evident that the oxygen gas, delivered to the O2 compartment, permeates
through the electrode membrane assembly and reacts with the water molecules present in the aqueous electrolyte and electrons
delivered from the electrode, producing further hydroxyl ions. Hence, the oxygen electrode becomes the positive electrode, or
cathode, where reduction occurs. Thus, an equilibrium process of hydroxyl ion consumption at the anode, and production on the
cathode, is established. Overall, there is no net loss or gain of hydroxyl ions; however, it is clear that the charge transfer within the
fuel cell is facilitated by this negative ion.
With reference to the electrochemical potentials in Table 2, the half cell reactions for the alkaline hydrogen fuel cell are
2H2 þ 4OH− ⇔4H2 O þ 4e−

−0:83 V

½24Š

O2 þ 2H2 O þ 4e− ⇔4OH−

þ0:4 V

½25Š

which gives an overall reaction for an AFC of
−


−

−

−

2H2 þ O2 þ /
2H2 O þ 4O/H þ 4/e ⇔4/H2 O þ 4e/ þ 4OH
/
2H2 þ O2 ⇔2H2 O
1
H2 þ O2 ⇔H2 O
2

½26Š

From eqns [24] and [25], the corresponding theoretical terminal voltage for an AFC is = 0.4 V – (–0.83 V) = 1.23 V, at STP.
As per eqn [19], this corresponds to a ‘free energy’ of
−ΔG ¼ Wuseful ¼ nFE
¼ 2Ã 96485Ã 1:23 ¼ 237353 Joules ¼ 237:4 kJ mol − 1
ΔG ¼ −237:4 kJ mol − 1

½27Š

However, the actual theoretical available energy from the reaction given in eqn [26] can be calculated from the ‘enthalpy of
combustion’ (which states the energy released from complete combustion of hydrogen in the presence of oxygen to form water (as a
liquid, i.e., at temperatures below 100 °C), as per eqn [26]) which is ΔH = –286 kJ mol−8. The recognition that the value of ΔH for
hydrogen is greater than that of the free energy, ΔG, indicates that not all the energy from the ‘reactants’ (i.e., the hydrogen and



Preface and Context to Hydrogen and Fuel Cells

25

oxygen gases) produces electrical energy. From the same thermodynamic data tables, for the reaction stated in eqn [26], the
corresponding values of enthalpy and entropy are as follows:
¼ 130:7 J ðK mol Þ − 1
¼ 205:0 J ðK mol Þ − 1
¼ 70 J ðK molÞ − 1
¼ −285:8 kJ mol − 1
�
� �
�
�
ΔS ¼ Sproducts – Sreactants ¼ SH2 O ðl Þ – SH2 ðg Þ þ 1 2 SO2 ðl Þ
�
¼ð70Þ–ð130:7 þ 1 2 à 205:0Þ ¼ 70 –233:2 ¼ −163:2 J ðK mol Þ − 1
SH2 ðg Þ
SO2 ðg Þ
SH2 O ðl Þ
ΔHH2 O ðl Þ

TΔS ¼ 298Ã ð−163:2Þ ¼ −48 633:6 J mol − 1

½28Š

ΔH ¼ ΔHH2 O ðl Þ ¼ −285:8 kJ mol − 1
Hence, the theoretical free energy, ΔG, is
ΔG ¼ ΔH−TΔS ¼ −285800–ð−48633:6Þ ¼ −237166:4 ¼ −237:2 kJ mol − 1

which is comparable with the value from eqn [27]. Equation [28] now indicates that theoretically a fair proportion of the available
energy from the hydrogen and oxygen fuel gases is released as heat: 48 633.6/285 800 = 17%.
It is worth noting that the product from the hydrogen fuel cell, that is, water, can be present in either liquid or gaseous form,
dependent on the operating temperature of the fuel cell. In the case of the AFC, the operating temperature is ∼70 °C which implies
that the water will be present in liquid form. For higher temperature fuel cells, such as SOFCs, which tend to operate at temperatures
from 600 to 1000 °C, the water product would be present in a gaseous state, that is, steam, and an alternative value for the enthalpy
and entropy of water would change to
SH2 O ðg Þ ¼ 188:8 J ðK mol Þ − 1
ΔHH2 O ðg Þ ¼ −241:8 kJ mol − 1
Hence, for the case of an SOFC at 1000 °C, the corresponding heat loss can be calculated as
�
� �
�
�
ΔS ¼ Sproducts – Sreactants ¼ SH2 O ðl Þ – SH2 ðg Þ þ 1 2 SO2 ðl Þ
�
¼ð188:8Þ –ð130:7 þ 1 2 à 205:0Þ ¼ 188:8 –233:2 ¼ −44:4 J ðK mol Þ − 1
TΔS ¼ 1273Ã ð−44:4Þ ¼ −56521 J mol − 1

½29Š

ΔG ¼ ΔH−TΔS ¼ −241 800–ð−56 521Þ ¼ −185279 ¼ −185:3 kJ mol − 1

½30Š

For this case, the resulting free energy is

This gives rise to a theoretical open-circuit cell voltage of, by rearranging eqns [18] and [19]
−ΔG ¼ nFE
−ΔG −ð−185300Þ

E¼
¼
¼ 0:960 V
nF
2Ã 96485

½31Š

This helps illustrate that as the working temperature of the fuel cell increases, the theoretical open-circuit voltage decreases
(compared to the voltage of 1.23 V at a temperature of 25 °C, from eqn [27]) and the ratio of energy ‘lost’ as waste heat increases.

4.01.4.4

Practical Fuel Cells

The theoretical open-circuit voltage from a hydrogen/oxygen fuel cell was considered and defined as 1.23 V at 25 °C. However, from
a practical perspective, a user would not be able to measure this voltage but rather measure a voltage typically < 1.1 V on open circuit
and, as an increasing electrical load is applied to the fuel cell, the terminal voltage decreases further. Figure 32 illustrates the typical
voltage versus current density (proportional to load current) characteristic for a fuel cell.
From Figure 32, it is clear that there are three distinct ‘regions’ to the fuel cell V/I characteristic [19]:
1. Activation polarization. This region experiences the initial energy loss required to activate the reactants and initiate the anode and
cathode half-cell reactions. This is evident at low currents and is nonlinear.
2. Ohmic polarization. This region exhibits the expected linear V/I characteristic apparent from electrically loading a nonideal voltage
source. This linear drop is caused by the voltage drop caused from a finite internal resistance (created by the resistances of the
electrolyte, electrode materials, electrical connections, and so forth).
3. Concentration polarization. This region is apparent only at high current densities (i.e., at or beyond full load) and is caused by the
particular cell design reaching its limit in terms of its ability to supply sufficient reactants to the cell and remove the products. In
essence, the cell cannot support the high level of reaction rate and effectively ‘throttles’ itself.